Periodic TrendsPeriodic trends are certain patterns that describe specific aspects of the elements in the periodic table, such as sizeand properties with electrons. The main periodic trends include: electronegativity, ionization energy, electronaffinity, atomic radius, melting point, and metallic character. The periodic trends that arise from the arrangementof the periodic table provide chemists with an invaluable tool to quickly predict an element's properties. Thesetrends exist because of the similar atomic structure of the elements within their respective group families or periodand the periodic nature of the elements.

Periodic Trends for Electronegativity . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1

Periodic Trends for Ionization Energy . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3

Periodic Trends for Electron Affinity . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 5

Periodic Trends for Atomic Radius . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 6

Periodic Trends for Melting Point. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 7

Periodic Trends for Metallic Character. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8

Summary. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 9

Outside Links . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 10

Problems . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 10

Solutions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 11

References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12

Contributors . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 13

Periodic Trends for ElectronegativityElectronegativity is a chemical property that describes an atom's ability of attracting and binding to electrons. Sinceelectronegativity describes a largely qualitative property, there is no standardized method of calculatingelectronegativity. However, the scale that most chemists use in quantifying electronegativity is the Pauling Scale,named after the chemist Linus Pauling. The numbers assigned by the Pauling scale are dimensionless due toelectronegativity being largely qualitative. Because electronegativity cannot be calculated, for the Chemistry 2series, students will be given a periodic table with electronegativity values for each element to work and study from.An example is provided below.

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Figure 1. Periodic Table of Electronegativity values

Electronegativity measures an atom's strength to attract and form bonds with electrons. This property exists due tothe electronic configuration of atoms. Most atoms prefer to fulfilling the octet rule (having the valence, or outer,shell comprise of 8 electrons). Since elements on the left side of the periodic table have less than a half-full valenceshell, the energy required to gain electrons is significantly higher compared to the energy required to lose electrons.As a result, the elements on the left side of the periodic table generally lose electrons in forming bonds. Conversely,elements on the right side of the periodic table are more energy-efficient in gaining electrons to create a completevalence shell of 8 electrons. This effectively describes the nature of electronegativity: the more inclined an atom isto gain electrons, the more likely that atom will pull electrons toward itself.

• As you move to the right across a period of elements, electronegativity increases. When the valence shell ofan atom is less than half full, it requires less energy to lose an electron than gain one and thus, it is easier tolose an electron. Conversely, when the valence shell is more than half full, it is easier to pull an electroninto the valence shell than to donate one.

• As you move down a group, electronegativity decreases. This is because the atomic number increases downa group and thus there is an increased distance between the valence electrons and nucleus, or a greateratomic radius.

• Important exceptions of the above rules include the noble gases, lanthanides, and actinides. The noble gasespossess a complete valence shell and do not usually attract electrons. The lanthanides and actinides possessa more complicated chemistry that does not generally follow any trends. Therefore, noble gases,lanthanides, and actinides do not have electronegativity values.

• As for the transition metals, while they have values, there is little variance among them as you move acrossthe period and up and down a group. This is because of their metaliic properties that affect their ability toattract electrons as easily as the other elements.

With these two general trends in mind, we can deduce that the most electronegative element is fluorine, which weighsin at a hefty 3.98 Pauling units.

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Figure 2. Periodic Table showing Electronegativity Trend

Periodic Trends for Ionization EnergyIonization Energy is the amount of energy required to remove an electron from a neutral atom in its gaseous phase.Conceptually, ionization energy is considered the opposite of electronegativity. The lower this energy is, the morereadily the atom becomes a cation. Therefore, the higher this energy is, the more unlikely the atom becomes acation. Generally, elements on the right side of the periodic table have a higher ionization energy because theirvalence shell is nearly filled. Elements on the left side of the periodic table have low ionization energies because oftheir willingness to lose electrons and become cations. Thus, ionization energy increases from left to right on theperiodic table.

Another factor that affects ionization energy is electron shielding. Electron shielding describes the ability of an atom'sinner electrons to shield its positively-charged nucleus from its valence electrons. When moving to the right on aperiod of elements, the number of electrons increases and the strength of shielding increases. As a result, it is easierfor valence shell electrons to ionize and thus the ionization energy decreases when going down a group. In certaintexts, electron shielding may also be known as screening.

• The ionization energy of the elements within a period generally increases from left to right. This is due tovalence shell stability.

• The ionization energy of the elements within a group generally decreases from top to bottom. This is due toelectron shielding.

• The noble gases possess very high ionization energies because of their full valence shell as indicated in thegraph. Note that Helium has the highest ionization energy of all the elements.

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Figure 3. Graph showing the Ionization Energy of the Elements from Hydrogen to Argon

Some elements can have several ionization energies, so we refer to these varying energies as the first ionizationenergy, the second ionization energy, third ionization energy, etc. The first ionization energy is to the energyneeded to remove the outermost, or highest, energy electron and the second ionization energy is the energyrequired to remove any subsequent high-energy electron from a gaseous cation. Below are the formulas forcalculating the first and second ionization energies.

First Ionization Energy:

Second Ionization Energy:

Generally, any subsequent ionization energies (2nd, 3rd, etc.) follow the same periodic trend as the first ionizationenergy.

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Figure 4. Periodic Table Showing Ionization Energy Trend

Ionization energies decrease as atomic radii increase. This observation is affected by n (the principle quantumnumber) and Zeff (based on the atomic number and shows how many protons are seen in the atom) on the ionizationenergy (I). Given by the following equation:

Going across a period, the Zeff increases and n (principal quantum number) remains the same, so that the ionizationenergy increases.

Going down a group, the n increases and Zeff increases slightly, the ionization energy decreases.

Periodic Trends for Electron AffinityLike the name suggests, electron affinity describes the ability of an atom to accept an electron. Unlikeelectronegativity, electron affinity is a quantitative measure that measures the energy change that occurs when anelectron is added to a neutral gas atom. When measuring electron affinity, the more negative the value, the more ofan affinity to electrons that atom has.

Electron affinity generally decreases down a group of elements because each atom is larger than the atom above it(this is the atomic radius trend, which will be discussed later in this text). This means that an added electron isfurther away from the atom's nucleus compared to its position in the smaller atom. With a larger distance betweenthe negatively-charged electron and the positively-charged nucleus, the force of attraction is relatively weaker.Therefore, electron affinity decreases. Moving from left to right across a period, atoms become smaller as the forcesof attraction become stronger. This causes the electron to move closer to the nucleus, thus increasing the electronaffinity from left to right across a period.

• Electron affinity increases from left to right within a period. This is caused by the decrease in atomic radius.• Electron affinity decreases from top to bottom within a group. This is caused by the increase in atomic radius.

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Figure 5. Periodic Table showing Electron Affinity Trend

Periodic Trends for Atomic RadiusFor atoms, the atomic radius is one-half the distance between the nuclei of two atoms is (just like a radius is half thediameter of a circle). However, this idea is complicated by the fact that not all atoms are normally bound together inthe same way. Some are bound by covalent bonds in molecules, some are attracted to each other in ionic crystals,and others are held in metallic crystals. Nevertheless, it is possible for a vast majority of elements to form covalentmolecules in which two like atoms are held together by a single covalent bond. The covalent radius of thesemolecules is often referred to as the atomic radius. This distance is measured in picometers. Going through each ofthe elements of the periodic table, patterns of the atomic radius can be seen.

Atomic size gradually decreases from left to right across a period of elements. This is because, within a period orfamily of elements, all electrons are being added to the same shell. But, at the same time, protons are being added tothe nucleus, making it more positively charged. The effect of increasing proton number is greater than that of theincreasing electron number; therefore, there is a greater nuclear attraction. This means that the nucleus attracts theelectrons more strongly, having the atom's shell pulled closer to the nucleus. The valence electrons are held closertowards the nucleus of the atom. As a result, the atomic radius decreases.

Going down a group, it can be seen that atomic radius increases. The valence electrons occupy higher levels due tothe increasing quantum number (n). As a result, the valence electrons are further away from the nucleus as the ‘n’increases. Electron shielding prevents these outer electrons from being attracted to the nucleus; thus, they areloosely held and the resulting atomic radius is large.

• Atomic radius decreases from left to right within a period. This is caused by the increase in the number ofprotons and electrons across a period. One proton has a greater effect than one electron; thus, a lot ofelectrons will get pulled towards the nucleus, resulting in a smaller radius.

• Atomic radius increases from top to bottom within a group. This is caused by electron shielding.

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Figure 6. Periodic Table showing Atomic Radius Trend

Periodic Trends for Melting PointMelting points are the amount of energy required to break a bond(s) to change the solid phase of a substance to aliquid. Generally, the stronger the bonds between the atoms of an element, the higher the energy requirement inbreaking that bond. Since temperature is directly proportional to energy, a high bond dissociation energy correlatesto a high temperature. Melting points are varied and don't generally form a distinguishable trend across theperiodic table. However, certain conclusions can be drawn from the following graph.

• Metals generally possess a high melting point.• Most non-metals possess low melting points.• The non-metal carbon possesses the highest boiling point of all the elements. The semi-metal boron also

possesses a high melting point.

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Figure 7. Chart of Melting Points of Various Elements

Periodic Trends for Metallic CharacterHow easily an atom can lose an electron is a measure of an element's metallic character. As you move from right toleft across a period, metallic character increases because the attraction between valence electron and the nucleus isweaker, thus enabling an easier loss of electrons. Metallic character increases as you move down a group because theatomic size is increasing . When the atomic size increases, the outer shells are farther away. The principle quantumnumber increases and average electron density moves farther from nucleus. The electrons of the valence shell haveless of an attraction to the nucleus and, as a result, can lose electrons more readily, causing an increase in metalliccharacter.

• Metallic characteristics decrease from left to right across a period. This is caused by the decrease in radius(above it is stated that Zeff causes this)of the atom which allows the outer electrons to ionize more readily.

• Metallic characteristics increase down a group. Electron shielding causes the atomic radius to increase thusthe outer electrons ionizes more readily than electrons in smaller atoms.

• Metallic character relates to the ability to lose electrons, and nonmetallic character relates to the ability togain electrons.

• Another easier way to remember the trend of metallic character is that as you move from left and downtowards the bottom-left corner of the periodic table, metallic character increases because you are headingtowards Groups 1 and 2, or the Alkali and Alkaline metal groups. Likewise, if you move up and to the rightto the upper-right corner of the periodic table, metallic character decreases because you are passing by tothe right side of the staircase, which indicate the nonmetals. These include the Group 8, the noble gases,and other common gases such as oxygen and nitrogen.

? In other words:? Move left across period and down the group: increase metallic character (heading towards alkali

and alkaline metals)? Move right across period and up the group: decrease metallic character (heading towards

nonmetals like noble gases)

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Figure 8. Periodic Table of Metallic Character Trend

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Summary

Outside Links• http://en.wikipedia.org/wiki/Periodic_trends• Pinto, Gabriel. "Using Balls of Different Sports To Model the Variation of Atomic Sizes." J. Chem. Educ.1998 75

725.• Qureshi, Pushkin M.; Kamoonpuri, S. Iqbal M. "Ion solvation: The ionic radii problem." J. Chem. Educ.1991, 68,

109.• Smith, Derek W. "Atomization enthalpies of metallic elemental substances using the semi-quantitative

theory of ionic solids: A simple model for rationalizing periodic trends." J. Chem. Educ. 1993, 70, 368.• Ionization Energy Trend: http://www.youtube.com/watch?v=ywqg9PorTAw&feature=youtube_gdata• Other Periodic Table Trends: http://www.youtube.com/watch?v=XMLd-O6PgVs&feature=youtube_gdata• Dr. Enderle's (UCD) Lecture on Periodic Trends (2 parts):

? (Part 1) http://www.youtube.com/watch?v=ZL1uXS43oQY&feature=relmfu? (Part 2) http://www.youtube.com/watch?v=p5YeD7jkz30&feature=relmfu

ProblemsThe following series of problems will review your general understanding of the aforementioned material.

1.) Based on the periodic trends for ionization energy, which do you except to have the highest ionization energy?

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1. A.) Fluorine (F)2. B.) Nitrogen (N)3. C.) Helium (He)

2.) Nitrogen has a larger atomic radius than Oxygen.

1. A.) True2. B.) False

3.) Which do you expect to have more metallic character, Lead (Pb) or Tin (Sn)?

4.) Which element do you expect to have the higher melting point: chlorine (Cl) or bromine (Br)?

5.) Which element do you expect to be more electronegative, sulfur (S) or selenium (Se)?

6) Why is the electronegativity value of most noble gases equal to zero?

7) Arrange the following atoms according to decreasing effective nuclear charge experienced by their valenceelectrons: S, Mg, Al, Si

8) Rewrite the following list in order of decreasing electron affinity: Fluorine (F), Phosphorous (P), Sulfur (S), Boron(B).

9) An atom with an atomic radius smaller than that of Sulfur (S) is __________.

1. A.) Oxygen (O)2. B.) Chlorine (Cl)3. C.) Calcium (Ca)4. D.) Lithium (Li)5. E.) None of the above

10) A nonmetal will have a smaller ionic radius when compared to a metal of the same period.

1. A.) True B.) False

11) Which one of the following has the lowest first ionization energy?

1. A. Element A2. B. Element B3. C. Element C4. D. Element D

Solutions1. Answer: C.) Helium (He)

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Explanation: Helium (He) has the highest ionization energy because, like other noble gases, Helium's valence shell isfull. Because of this, Helium is stable and does not readily lose or gain electrons.

2. Answer: A.) True

Explanation: According to periodic trends, atomic radius increases from right to left on the periodic table.Therefore, we would expect Nitrogen to be larger than Oxygen.

3. Answer: Lead (Pb)

Explanation: Lead and Tin share the same column. According to periodic trends, metallic character increases as yougo down a column. Lead is underneath Tin therefore we would expect Lead to possess more metallic character.

4. Answer: Bromine (Br)

Explanation: According to periodic trends, in non-metals, melting point increases down a column. Since chlorineand bromine share the same column, we would expect bromine to possess the higher melting point.

5. Answer: Sulfur (S)

Explanation: Note that sulfur and selenium share the same column. Periodic trends tell us that electronegativityincreases up a column. This indicates that sulfur is more electronegative than selenium.

6. Answer: Most noble gases have full valence shells.

Explanation: Because of their full valence electron shell, the noble gases are extremely stable and do not readily loseor gain electrons.

7. Answer: S > Si > Al > Mg.

Explanation: The electrons above a closed shell are shielded by the closed shell. S has 6 electrons above a closedshell, so each one feels the pull of 6 prontons in the nucleus.

8. Answer: Fluorine (F)>Sulfur (S)>Phosphorous (P)>Boron (B)

Explanation: According to periodic trends, the electron affinity generally increases from left to right and frombottom to top.

9. Answer: C.) Oxygen (O)

Explanation: Periodic trends indicate that atomic radius increases up a group and from left to right across a period.Therefore, oxygen is expect to have a smaller atomic radius than of sulfur.

10. Answer: B.) False

Explanation: The reasoning behind this lies in understanding that a metal usually loses an electron in becoming anion while a non-metal gains an electron. This results in a smaller ionic radius for the metal ion and a larger ionicradius for the non-metal ion.

11. Element D

Explanation: Element A, B and D have the same number of elentrons in the inner shell, but element D has the leastnumber of eletrons in the outer shell which requires the lowest ionization energy.

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References1. Russo, Steve, and Mike Silver. Introductory Chemistry. San Francisco: Pearson, 2007.2. Petrucci, Ralph H, et al. General Chemistry: Principles and Modern Applications. 9th Ed. New Jersey:

Pearson, 2007.3. Atkins, Peter et. al, Physical Chemistry, 7th Edition, 2002, W.H Freeman and Company, New York, pg. 390.4. Alberty, Robert A. et. al, Physical Chemistry, 3rd Edition, 2001, John Wiley & Sons, Inc, pg. 380.5. Kots, John C. et. al, Chemistry & Chemical Reactivity, 5th Edition, 2003, Thomson Learning Inc, pg. 305-309.

Contributors• Swetha Ramireddy (UCD)• Bingyao Zheng (UCD)• Emily Nguyen (UCD)

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